Published on

Published in: Business
  • Be the first to comment

No Downloads
Total views
On SlideShare
From Embeds
Number of Embeds
Embeds 0
No embeds

No notes for slide


  1. 1. Chapter 2 Gases
  2. 2. 12.1 Characteristics of Gases <ul><li>Properties of Gases </li></ul><ul><ul><li>because gas particles are far apart, </li></ul></ul><ul><ul><li>gases are fluids (they can flow) </li></ul></ul><ul><ul><li>gases have low density </li></ul></ul><ul><ul><li>gases are highly compressible </li></ul></ul><ul><ul><li>gases completely fill a container </li></ul></ul>
  3. 3. 12.1 Characteristics of Gases <ul><li>Gas Pressure </li></ul><ul><ul><li>Rene Descarte (1596-1650): rejected idea of void or vacuum </li></ul></ul><ul><ul><li>Pierre Gassendi (1592-1655): revived atomism; promoted idea of atoms moving in a void </li></ul></ul><ul><ul><li>Evangelista Torricelli (1608-1647): built a mercury barometer in 1643; created a vacuum </li></ul></ul>
  4. 4. Mercury Barometer
  5. 5. 12.1 Characteristics of Gases <ul><li>Gas Pressure </li></ul><ul><ul><li>Blaise Pascal (1623-1662): tested atmospheric pressure at prompting of Descarte; found that pressure drops with altitude; believed in the vacuum </li></ul></ul>
  6. 6. 12.1 Characteristics of Gases <ul><li>Gas Pressure </li></ul><ul><ul><li>pressure is force divided by area </li></ul></ul><ul><ul><ul><li>force: Newton (1 kg  m/s 2 = 1 N) </li></ul></ul></ul><ul><ul><ul><li>area: meter squared (m 2 ) </li></ul></ul></ul><ul><ul><ul><li>pressure: Pascal (1 Pa = 1 N/1 m 2 ) </li></ul></ul></ul><ul><ul><li>for comparisons, standard temperature and pressure (STP): 0  C and 1 atm </li></ul></ul>
  7. 7. Pressure Units
  8. 8. 12.1 Characteristics of Gases <ul><li>Kinetic-Molecular Theory </li></ul><ul><ul><li>gas particles are in constant, rapid, random motion </li></ul></ul><ul><ul><li>particles far apart relative to size </li></ul></ul><ul><ul><li>pressure due to collisions of particles with the walls of their container </li></ul></ul>
  9. 9. 12.1 Characteristics of Gases <ul><li>Kinetic-Molecular Theory </li></ul><ul><ul><li>gas temperature is proportional to average kinetic energy </li></ul></ul><ul><ul><ul><li>gas molecules have a range of speeds </li></ul></ul></ul><ul><ul><ul><li>increasing temperature shifts the distribution </li></ul></ul></ul>
  10. 10. Gas Molecules Energy Distribution
  11. 11. 12.2 The Gas Laws <ul><li>Measurable Properties of Gases </li></ul><ul><ul><li>P = pressure exerted by gas </li></ul></ul><ul><ul><li>V = total volume occupied by gas </li></ul></ul><ul><ul><li>T = temperature in kelvins of gas </li></ul></ul><ul><ul><li>n = number of moles of gas </li></ul></ul>
  12. 12. 12.2 The Gas Laws <ul><li>Robert Boyle (1627-1691): published The Spring of Air in 1660, which explained his most famous experiment </li></ul><ul><ul><li>Boyle put mercury in a j-tube (manometer), and saw that when he doubled the pressure, the volume of air in short end halved </li></ul></ul>
  13. 13. Boyle’s Experiment
  14. 14. Boyle’s Law
  15. 15. 12.2 The Gas Laws <ul><li>Robert Boyle </li></ul><ul><ul><li>Boyle’s law: </li></ul></ul><ul><ul><ul><li>PV = k </li></ul></ul></ul><ul><ul><ul><li>P 1 V 1 = P 2 V 2 </li></ul></ul></ul>
  16. 16. Boyle’s Law
  17. 17. 12.2 The Gas Laws <ul><li>Jacques Charles: discovered that a gas’s volume is proportional to temperature at constant pressure in 1787 </li></ul><ul><ul><li>Charles’s law: </li></ul></ul><ul><ul><ul><li>V / T = k </li></ul></ul></ul><ul><ul><ul><li>V 1 / T 1 = V 2 / T 2 </li></ul></ul></ul>
  18. 18. Charles’s Law
  19. 19. 12.2 The Gas Laws <ul><li>Joseph Gay-Lussac (1778-1850): discovered in 1802 that increasing temperature at constant volume resulted in a proportional increase in pressure </li></ul><ul><ul><li>Gay-Lussac’s law: </li></ul></ul><ul><ul><ul><li>P = kT </li></ul></ul></ul><ul><ul><ul><li>P / T = k </li></ul></ul></ul><ul><ul><ul><li>P 1 / T 1 = P 2 / T 2 </li></ul></ul></ul>
  20. 20. Gay-Lussac’s Law
  21. 21. 12.2 The Gas Laws <ul><ul><li>Gay-Lussac’s law of combining volumes (1809): gases combine in simple proportions by volume, and volume of products is related to volume of reactants </li></ul></ul><ul><ul><ul><li>example 1: 2 volumes of H 2 react with 1 volume of O 2 to make 2 volumes of water </li></ul></ul></ul><ul><ul><ul><li>allowed Avogadro to deduce diatomic molecules (and more) </li></ul></ul></ul>
  22. 22. Combining Volumes
  23. 23. 12.2 The Gas Laws <ul><li>Amadeo Avogadro (1776-1856): proposed in 1811 that equal volumes of all gases contain equal numbers of particles </li></ul><ul><ul><li>Avogadro’s law: </li></ul></ul><ul><ul><ul><li>V = kn </li></ul></ul></ul><ul><ul><ul><li>1 mol of any gas at 0  C and 1 atm occupies 22.41 L </li></ul></ul></ul>
  24. 24. Avogadro’s Law
  25. 25. 12.2 The Gas Laws <ul><li>Stanislao Cannizzaro (1826-1910): ~1858, deduced that Gay-Lussac’s law of combining volumes and Avogadro’s law could be used to calculate atomic and molecular weights relative to hydrogen; drew distinction between atoms and molecules; made a table of atomic weights </li></ul>
  26. 26. Gas Laws Summary
  27. 27. 12.3 Molecular Comp. of Gases <ul><li>Ideal Gas Law </li></ul><ul><ul><li>no gas perfectly obeys Boyle’s law, Charles’s law, Gay-Lussac’s law, or Avogadro’s law </li></ul></ul><ul><ul><li>although not perfect, these laws work well for most gases and most conditions </li></ul></ul><ul><ul><li>ideal gas : model gas that perfectly obeys gas laws </li></ul></ul>
  28. 28. Ideal Gases vs. Real Gases
  29. 29. 12.3 Molecular Comp. of Gases <ul><li>Ideal Gas Law </li></ul><ul><ul><li>ideal gases </li></ul></ul><ul><ul><ul><li>do not condense to liquids at low temperatures </li></ul></ul></ul><ul><ul><ul><li>do not have particles attracted to or repulsed by each other </li></ul></ul></ul><ul><ul><ul><li>have particles of no volume </li></ul></ul></ul><ul><ul><ul><li>do not exist </li></ul></ul></ul>
  30. 30. 12.3 Molecular Comp. of Gases <ul><li>Ideal Gas Law : combines four variables, P , V , T , and n , into one equation </li></ul><ul><ul><li>PV = nRT </li></ul></ul><ul><ul><li>R is a proportionality constant </li></ul></ul><ul><ul><li>R = 8.314 L  kPa </li></ul></ul><ul><ul><li> mol  K </li></ul></ul>
  31. 31. 12.3 Molecular Comp. of Gases <ul><li>Gas Behavior and Chemical Formulas </li></ul><ul><ul><li>Diffusion : movement of particles from high concentration to low concentration </li></ul></ul><ul><ul><ul><li>particles of lower mass diffuse more quickly than particles of higher mass </li></ul></ul></ul><ul><ul><ul><li>diffusion increases entropy </li></ul></ul></ul>
  32. 32. 12.3 Molecular Comp. of Gases <ul><li>Gas Behavior and Chemical Formulas </li></ul><ul><ul><li>Effusion : passage of gas particles through a small opening </li></ul></ul><ul><ul><ul><li>Graham’s law: rate of diffusion and effusion of a gas are inversely proportional to the square root of the gas’s density </li></ul></ul></ul>
  33. 33. 12.3 Molecular Comp. of Gases <ul><li>Gas Behavior and Chemical Formulas </li></ul><ul><ul><ul><li>Graham’s law, cont. </li></ul></ul></ul><ul><ul><ul><ul><li>where v A and v B are molecular speeds of gases A and B and </li></ul></ul></ul></ul><ul><ul><ul><ul><li>M A and M B are the molar masses of gases A and B </li></ul></ul></ul></ul>
  34. 34. 12.3 Molecular Comp. of Gases <ul><li>Gas Behavior and Chemical Formulas </li></ul><ul><ul><ul><li>Graham’s law, cont. </li></ul></ul></ul><ul><ul><ul><ul><li>Graham’s law is easy to derive: solve the equation for the ratio of speeds between v A and v B </li></ul></ul></ul></ul>
  35. 35. 12.3 The Gas Laws <ul><li>John Dalton (1766-1844): discovered that each gas in a mixture produces its own pressure as if it was alone </li></ul><ul><ul><li>Dalton’s law of partial pressure: total pressure of a mixture of gases is equal to the sum of the partial pressures of the component gases </li></ul></ul><ul><ul><ul><li>P total = P A + P B + P C </li></ul></ul></ul>